Calcium carbonate

 ( Calcium Compounds ) 

At high saturation, vaterite is typically the first phase precipitated, which is followed by a transformation of the vaterite to calcite. This behavior seems to follow Ostwald's rule, in which the least stable polymorph crystallizes first, followed by the crystallization of different polymorphs via a sequence of increasingly stable phases. However, aragonite, whose stability lies between those of vaterite and calcite, seems to be the exception to this rule, as aragonite does not form as a precursor to calcite under ambient conditions. Aragonite occurs in majority when the reaction conditions inhibit the formation of calcite and/or promote the nucleation of aragonite. For example, the formation of aragonite is promoted by the presence of magnesium ions, or by using proteins and peptides derived from biological calcium carbonate. Some polyamines such as cadaverine and Polyethylenimine have been shown to facilitate the formation of aragonite over calcite. Solid‑state NMR analysis has revealed that poly‑aspartate-stabilized ACC contains water molecules that undergo millisecond-timescale flips, illustrating dynamic hydration as a key factor in delaying crystallization.

in the family Lumbricidae, earthworms, possess a regionalization of the digestive track called calciferous glands, Kalkdrüsen, or glandes de Morren, that processes calcium and Carbon dioxide into calcium carbonate, which is later excreted into the dirt. The function of these glands is unknown but is believed to serve as a regulation mechanism within the animals' tissues. This process is ecologically significant, stabilizing the Soil pH.

Where the oceanic crust is Subduction under a continental plate sediments will be carried down to warmer zones in the asthenosphere and lithosphere. Under these conditions calcium carbonate decomposes to produce carbon dioxide which, along with other gases, give rise to explosive volcano.

Calcium carbonate is also used in the purification of iron from iron ore in a blast furnace. The carbonate is Calcination in situ to give calcium oxide, which forms a slag with various impurities present, and separates from the purified iron.

In the oil industry, calcium carbonate is added to as a formation-bridging and filtercake-sealing agent; it is also a weighting material which increases the density of drilling fluids to control the downhole pressure. Calcium carbonate is added to , as a pH corrector for maintaining alkalinity and offsetting the acidic properties of the disinfectant agent.

It is also used as a raw material in the Sugar refining from sugar beet; it is calcined in a kiln with anthracite to produce calcium oxide and carbon dioxide. This burnt lime is then slaked in fresh water to produce a calcium hydroxide suspension for the precipitation of impurities in raw juice during carbonatation.

Calcium carbonate in the form of chalk has traditionally been a major component of blackboard chalk. However, modern manufactured chalk is mostly gypsum, hydrated calcium sulfate . Calcium carbonate is a main source for growing biorock. Precipitated calcium carbonate (PCC), pre-dispersed in slurry form, is a common filler material for latex gloves with the aim of achieving maximum saving in material and production costs.

Fine ground calcium carbonate (GCC) is an essential ingredient in the microporous film used in diapers and some building films, as the pores are nucleated around the calcium carbonate particles during the manufacture of the film by biaxial stretching. GCC and PCC are used as a filler in paper because they are cheaper than wood fiber. Printing and writing paper can contain 10–20% calcium carbonate. In North America, calcium carbonate has begun to replace Kaolinite in the production of glossy paper. Europe has been practicing this as alkaline papermaking or acid-free papermaking for some decades. PCC used for paper filling and paper coatings is precipitated and prepared in a variety of shapes and sizes having characteristic narrow particle size distributions and equivalent spherical diameters of 0.4 to 3 micrometers.

Calcium carbonate is widely used as an extender in , in particular matte emulsion paint where typically 30% by weight of the paint is either chalk or marble. It is also a popular filler in plastics. Some typical examples include around 15–20% loading of chalk in unplasticized polyvinyl chloride (uPVC) Rain gutter, 5–15% loading of stearic acid-coated chalk or marble in uPVC window profile. PVC cables can use calcium carbonate at loadings of up to 70 phr (parts per hundred parts of resin) to improve mechanical properties (tensile strength and elongation) and electrical properties (volume resistivity). Polypropylene compounds are often filled with calcium carbonate to increase rigidity, a requirement that becomes important at high usage temperatures. Here the percentage is often 20–40%. It also routinely used as a filler in thermosetting resins (sheet and bulk molding compounds) and has also been mixed with ABS, and other ingredients, to form some types of compression molded "clay" . Precipitated calcium carbonate, made by dropping calcium oxide into water, is used by itself or with additives as a white paint, known as .

Calcium carbonate is added to a wide range of trade and do it yourself adhesives, sealants, and decorating fillers. Ceramic tile adhesives typically contain 70% to 80% limestone. Decorating crack fillers contain similar levels of marble or dolomite. It is also mixed with putty in setting stained glass windows, and as a resist to prevent glass from sticking to kiln shelves when firing glazes and paints at high temperature.

In ceramic glaze applications, calcium carbonate is known as whiting, and is a common ingredient for many glazes in its white powdered form. When a glaze containing this material is fired in a kiln, the whiting acts as a Ceramic flux material in the glaze. Ground calcium carbonate is an abrasive (both as scouring powder and as an ingredient of household scouring creams), in particular in its calcite form, which has the relatively low hardness level of 3 on the Mohs scale, and will therefore not scratch glass and most other , Vitreous enamel, bronze, iron, and steel, and have a moderate effect on softer metals like aluminium and copper. A paste made from calcium carbonate and deionized water can be used to clean tarnish on silver.

Calcium carbonate is used in the production of calcium oxide as well as toothpaste and has seen a resurgence as a food preservative and color retainer, when used in or with products such as organic apples.

Calcium carbonate is used therapeutically as phosphate binder in patients on maintenance haemodialysis. It is the most common form of phosphate binder prescribed, particularly in non-dialysis chronic kidney disease. Calcium carbonate is the most commonly used phosphate binder, but clinicians are increasingly prescribing the more expensive, non-calcium-based phosphate binders, particularly sevelamer.

Excess calcium from supplements, fortified food, and high-calcium diets can cause milk-alkali syndrome, which has serious toxicity and can be fatal. In 1915, Bertram Sippy introduced the "Sippy regimen" of hourly ingestion of milk and cream, and the gradual addition of eggs and cooked cereal, for 10 days, combined with alkaline powders, which provided symptomatic relief for peptic ulcer disease. Over the next several decades, the Sippy regimen resulted in kidney failure, alkalosis, and hypercalcaemia, mostly in men with peptic ulcer disease. These adverse effects were reversed when the regimen stopped, but it was fatal in some patients with protracted vomiting. Milk-alkali syndrome declined in men after effective treatments for peptic ulcer disease arose. Since the 1990s it has been most frequently reported in women taking calcium supplements above the recommended range of 1.2 to 1.5 grams daily, for prevention and treatment of osteoporosis, and is exacerbated by dehydration. Calcium has been added to over-the-counter products, which contributes to inadvertent excessive intake. Excessive calcium intake can lead to hypercalcemia, complications of which include vomiting, abdominal pain and altered mental status.

As a food additive it is designated E numbers, 080419 food-info.net and it has an INS number of 170. Used as an acidity regulator, anticaking agent, stabilizer or Food coloring it is approved for usage in the EU, US and Australia and New Zealand. It is "added by law to all UK milled bread flour except wholemeal". It is used in some soy milk and almond milk products as a source of dietary calcium; at least one study suggests that calcium carbonate might be as bioavailable as the calcium in cow's milk. Calcium carbonate is also used as a firming agent in many canned and bottled vegetable products.

Several calcium supplement formulations have been documented to contain the chemical element lead, posing a public health concern. Lead is commonly found in natural sources of calcium.

Calcium carbonate is also used in flue-gas desulfurization applications eliminating harmful and emissions from coal and other fossil fuels burnt in large fossil fuel power stations.

At temperatures above 550 °C the equilibrium pressure begins to exceed the pressure in air. So above 550 °C, calcium carbonate begins to outgas into air. However, in a charcoal fired kiln, the concentration of will be much higher than it is in air. Indeed, if all the oxygen in the kiln is consumed in the fire, then the partial pressure of in the kiln can be as high as 20 kPa.

The table shows that this partial pressure is not achieved until the temperature is nearly 800 °C. For the outgassing of from calcium carbonate to happen at an economically useful rate, the equilibrium pressure must significantly exceed the ambient pressure of . And for it to happen rapidly, the equilibrium pressure must exceed total atmospheric pressure of 101 kPa, which happens at 898 °C.

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+ Equilibrium pressure of over ( P) versus temperature ( T).

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Solubility

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With varying pressure

Calcium carbonate is poorly soluble in pure water (47 mg/L at normal atmospheric partial pressure as shown below).

The equilibrium of its solution is given by the equation (with dissolved calcium carbonate on the right):

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Ksp = to at 25 °C

where the solubility product for is given as anywhere from K_sp = to K_sp = at 25 °C, depending upon the data source. What the equation means is that the product of molar concentration of calcium ions (moles of dissolved per liter of solution) with the molar concentration of dissolved cannot exceed the value of K_sp. This seemingly simple solubility equation, however, must be taken along with the more complicated equilibrium of carbon dioxide with water (see carbonic acid). Some of the combines with in the solution according to

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Ka2 = at 25 °C

is known as the [[bicarbonate]] ion. Calcium bicarbonate is many times more soluble in water than calcium carbonate—indeed it exists ''only'' in solution.
     

Some of the combines with in solution according to

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Ka1 = at 25 °C

Some of the breaks up into water and dissolved carbon dioxide according to

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Kh = at 25 °C

And dissolved carbon dioxide is in equilibrium with atmospheric carbon dioxide according to

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| width=250 = $H\_\{\backslash rm\; v\}$

where $H\_\{\backslash rm\; v\}$ = 29.76 atm/(mol/L) at 25 °C (Henry volatility), and P is the partial pressure.

For ambient air, P is around atm (or equivalently 35 Pa). The last equation above fixes the concentration of dissolved as a function of P, independent of the concentration of dissolved . At atmospheric partial pressure of , dissolved concentration is moles per liter. The equation before that fixes the concentration of as a function of concentration. For = , it results in = moles per liter. When is known, the remaining three equations together with

+ Calcium ion solubility as a function of carbon dioxide partial pressure at 25 °C ( Ksp = )

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K = 10−14 at 25 °C

(which is true for all aqueous solutions), and the constraint that the solution must be electrically neutral, i.e., the overall charge of dissolved positive ions must be cancelled out by the overall charge of dissolved negative ions , make it possible to solve simultaneously for the remaining five unknown concentrations (the previously mentioned form of the neutrality is valid only if calcium carbonate has been put in contact with pure water or with a neutral pH solution; in the case where the initial water solvent pH is not neutral, the balance is not neutral).

The adjacent table shows the result for and (in the form of pH) as a function of ambient partial pressure of ( K_sp = has been taken for the calculation).

• At atmospheric levels of ambient the table indicates that the solution will be slightly alkaline with a maximum solubility of 47 mg/L.
• As ambient partial pressure is reduced below atmospheric levels, the solution becomes more and more alkaline. At extremely low P, dissolved , bicarbonate ion, and carbonate ion largely evaporate from the solution, leaving a highly alkaline solution of calcium hydroxide, which is more soluble than . For P = 10⁻¹² atm, the product is still below the solubility product of (). For still lower pressure, precipitation will occur before precipitation.
• As ambient partial pressure increases to levels above atmospheric, pH drops, and much of the carbonate ion is converted to bicarbonate ion, which results in higher solubility of .

The effect of the latter is especially evident in day-to-day life of people who have hard water. Water in aquifers underground can be exposed to levels of much higher than atmospheric. As such, water percolates through calcium carbonate rock, the dissolves according to one of the trends above. When that same water then emerges from the tap, in time, it comes into equilibrium with levels in the air by outgassing its excess . The calcium carbonate becomes less soluble as a result, and the excess precipitates as lime scale. This same process is responsible for the formation of stalactites and in limestone caves.

Two hydrated phases of calcium carbonate, monohydrocalcite and ikaite , may precipitate from water at ambient conditions and persist as metastable phases.

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With varying pH, temperature and salinity: scaling in swimming pools

In contrast to the open equilibrium scenario above, many swimming pools are managed by addition of sodium bicarbonate () to the concentration of about 2 mmol/L as a buffer, then control of pH through use of HCl, , , NaOH or chlorine formulations that are acidic or basic. In this situation, dissolved inorganic carbon (total inorganic carbon) is far from equilibrium with atmospheric . Progress towards equilibrium through outgassing of is slowed by

In this situation, the dissociation constants for the much faster reactions

allow the prediction of concentrations of each dissolved inorganic carbon species in solution, from the added concentration of (which constitutes more than 90% of Bjerrum plot species from pH 7 to pH 8 at 25 °C in fresh water). Addition of will increase concentration at any pH. Rearranging the equations given above, we can see that = , and = . Therefore, when concentration is known, the maximum concentration of ions before scaling through precipitation can be predicted from the formula:

_max = ×

The solubility product for ( K_sp) and the dissociation constants for the dissolved inorganic carbon species (including K_a2) are all substantially affected by temperature and salinity, with the overall effect that _max increases from freshwater to saltwater, and decreases with rising temperature, pH, or added bicarbonate level, as illustrated in the accompanying graphs.

The trends are illustrative for pool management, but whether scaling occurs also depends on other factors including interactions with , and other ions in the pool, as well as supersaturation effects. Scaling is commonly observed in electrolytic chlorine generators, where there is a high pH near the cathode surface and scale deposition further increases temperature. This is one reason that some pool operators prefer borate over bicarbonate as the primary pH buffer, and avoid the use of pool chemicals containing calcium.

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Solubility in a strong or weak acid solution

Solutions of strong acid (HCl), moderately strong (sulfamic acid) or weak acid (acetic acid, citric acid, sorbic acid, lactic acid, phosphoric acid) acids are commercially available. They are commonly used as to remove limescale deposits. The maximum amount of that can be "dissolved" by one liter of an acid solution can be calculated using the above equilibrium equations.

• In the case of a strong monoacid with decreasing acid concentration A = , we obtain (with molar mass = 100 g/mol):

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! A (mol/L) 1	10−1	10−2	10−3	10−4	10−5	10−6	10−7	10−10

where the initial state is the acid solution with no (not taking into account possible dissolution) and the final state is the solution with saturated . For strong acid concentrations, all species have a negligible concentration in the final state with respect to and so that the neutrality equation reduces approximately to 2 = yielding ≈ 0.5 . When the concentration decreases, becomes non-negligible so that the preceding expression is no longer valid. For vanishing acid concentrations, one can recover the final pH and the solubility of in pure water.

• In the case of a weak monoacid (here we take acetic acid with p K_a = 4.76) with decreasing total acid concentration A = + AH, we obtain:

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!A (mol/L) ≈ 0.5	10−1	10−2	10−3	10−4	10−5	10−6	10−7	10−10

For the same total acid concentration, the initial pH of the weak acid is less acid than that of the strong acid; however, the maximum amount of which can be dissolved is approximately the same. This is because in the final state, the pH is larger than the p K_a, so that the weak acid is almost completely dissociated, yielding in the end as many ions as the strong acid to "dissolve" the calcium carbonate.

• The calculation in the case of phosphoric acid (which is the most widely used for domestic applications) is more complicated since the concentrations of the four dissociation states corresponding to this acid must be calculated together with , , , and . The system may be reduced to a seventh degree equation for the numerical solution of which gives

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! A (mol/L) 1	10−1	10−2	10−3	10−4	10−5	10−6	10−7	10−10

where A = is the total acid concentration. Thus phosphoric acid is more efficient than a monoacid since at the final almost neutral pH, the second dissociated state concentration is not negligible (see phosphoric acid).

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See also

• Cuttlebone
• Cuttlefish
• Gesso
• Limescale
• Marble
• Ocean acidification
• Whiting event
• List of climate engineering topics
• Lysocline

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External links

• ATC codes: and
• The British Calcium Carbonate Association – What is calcium carbonate
• CDC – NIOSH Pocket Guide to Chemical Hazards – Calcium Carbonate

Categories: Calcium Compounds, Carbonates, Limestone, Phosphate Binders, Excipients, Antacids, Food Stabilizers, E-number Additives, Over-the-counter Drugs In The United States